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CHAPTER 2 NOTES
2–1 The Nature of Matter
A. Atoms – the basic unit of matter- “unable to cut”
-made up of sub-atomic particles protons (+),neutrons (no chg), and electrons (-). Protons and neutrons are about the same mass, but electrons are much smaller - 1/1840 the mass of a proton.
Atomic Number= # of protons in an atom= # of electrons in an atom
Atomic Mass Number = the average number of protons and neutrons combined in an atom
B. Elements and Isotopes
1. Isotopes - Atoms of the same element that differ in the number of neutrons they contain. Because they have the same numbers of electrons, all isotopes of an element have the same chemical properties. Isotopes are identified by their mass numbers (C-14)
2. Radioactive Isotopes- radioactive, meaning that their nuclei are unstable and break down. The radiation given off can be dangerous, but it has many uses- to determine the ages of fossils by analyzing the isotopes found in them, and can be used to treat cancer. Radioactive isotopes (Ex. C-14,P-32) can be used as labels or “tracers” to follow the movements of substances within organisms. They are detectible by some analytical equipment (Ex. Geiger meter, scintillation counter)
C. Chemical Compounds- substances formed by the chemical combination of two or more elements in definite proportions.
Molecules are the smallest unit of compounds whose atoms are joined by covalent bonds
The physical and chemical properties of a compound are usually very different from those of the elements from which it is formed
Example:
Chemical formulas tell what elements are in a compound and the ratio in which the atoms combine
H2O is two hydrogen atoms and one oxygen atom in a ratio of 2:1, hydrogen to oxygen
D. Chemical Bonds-main types are ionic and covalent
Ions are positively and negatively charged atoms
1. Ionic Bonds- are formed when one or more electrons are transferred from one atom to another; the attraction of two oppositely charged ions
TWO STEPS INVOLVED:
· The transfer of electrons between atoms to fill valence levels
This step results in atoms with a + and a – charge
· The + and – come together and form an ionic bond
2. Covalent Bonds- electrons are shared between atoms. The outer energy levels (orbitals) overlap, and the shared electrons are in the orbitals of both atoms. Sharing one pair of electrons (2 electrons) is a single covalent bond- can have double and triple covalent bonds
· The type of bonds that hold molecules together
·
WHEN OUTER ENERGY (VALENCE) LEVELS ARE FILLED ARE FILLED, THE ATOM/ION IS STABLE AND UNREACTIVE (INERT)
2. Van der Waals Forces- intermolecular forces between oppositely charged regions of nearby molecules (much weaker than ionic or covalent bonds)
Ex.
E. Periodic Chart
1. Groups (columns) –(up/down) have same number of valence electrons
2. Left side – metallic elements that tend to LOSE electrons
3. Right side – non-metallic elements that tend to GAIN electrons
2–2 Properties of Water
A. The Water Molecule- neutral (proton # = electron #), BUT there is more to the story…..!
1. Polarity
A water molecule is polar because there is an uneven distribution of the shared electrons between the oxygen and hydrogen atoms. The negative pole is near the oxygen atom (it has 8 protons, and can attract the shared electrons more than the 1 proton in the H atom). The positive pole is between the hydrogen atoms.
3. Hydrogen Bonds
The attraction between the hydrogen atom on one water molecule and the oxygen atom on another water molecule is an example of a hydrogen bond. Hydrogen bonds are not as strong as covalent or ionic bonds, but they are stronger than Van der Waals forces.
Cohesion –“common sticking”, like a ‘chain’ of molecules- an attraction between molecules of the same substance. Because of hydrogen bonding, water is extremely cohesive. Ex. insects can walk on water’s surface
Adhesion- “toward different sticking” is an attraction between molecules of different substances. Ex. water rising up the sides of a glass tube (capillary action)
B. Solutions and Suspensions- A mixture is a material composed of two or more elements or compounds that are physically mixed together but not chemically combined. Ex. salt and pepper and all living things (which are mixed with water!)
1. Solutions- ions breaking away from their crystal form and becoming dispersed in the water, forming a type of mixture called a solution. solute—the substance that is dissolved. Water is the solvent—the substance in which the solute dissolves. Water’s polarity gives it the ability to dissolve both ionic compounds and other polar molecules
2. Suspensions - mixtures of water and nondissolved material are known as suspensions. Ex. Blood is a suspension of things like red/white blood cells-the small particles are suspended by movement of the water molecules
C. Acids, Bases, and pH
A water molecule can react to form ions. This reaction can be summarized by a chemical equation in which double arrows are used to show that the reaction can occur in either direction.
1. pH- “the potential of hydrogen”-
2. the pH Scale- the concentration of hydrogen ions in solution
A scale of 0-14. High values are basic, low values are acidic. At a pH of 7, the concentration of H+ ions and OH− ions is equal. Pure water has a pH of 7 (see equation above)
3. Acids - Acidic solutions contain higher concentrations of H+ ions than pure water and have pH values below 7, and are called acidic because they have more H+ ions than OH− ions.
4. Bases - Basic, or alkaline, solutions contain lower concentrations of H+ ions than pure water and have pH values above 7, and have more OH− ions than H+ ions
5. Buffers- Buffers are weak acids or bases that can react with strong acids or bases to prevent sharp, sudden changes in pH. Controlling pH is important for maintaining homeostasis.
2–3 Carbon Compounds
· Four groups of organic compounds found in living things are carbohydrates, lipids, nucleic acids, and proteins.
· Living things use carbohydrates as their main source of energy. Plants and some animals also use carbohydrates for structural purposes.
· Lipids can be used to store energy. Some lipids are important parts of biological membranes and waterproof coverings.
· Nucleic acids store and transmit hereditary, or genetic, information.
· Some proteins control the rate of reactions and regulate cell processes. Some proteins build tissues such as bones and muscles. Others transport materials or help to fight diseases.
A. The Chemistry of Carbon
B. Macromolecules
C. Carbohydrates
D. Lipids
E. Nucleic Acids
F. Protein
2–4 Chemical Reactions and Enzymes
· Chemical reactions always involve the breaking of bonds in reactants and the formation of new bonds in products.
· Chemical reactions that release energy often occur spontaneously. Chemical reactions that absorb energy will not occur without a source of energy.
· Enzymes speed up chemical reactions that take place in cells.
A. Chemical Reactions
B. Energy in Reactions
1. Energy Changes
2. Activation Energy
C. Enzymes
D. Enzyme Action
1. The Enzyme-Substrate Complex
2. Regulation of Enzyme Activity